1. SOL Review 3 Bonding and Naming

2. Chemical Bonding Ionic bonds occur between a metal and a nonmetal. Electrons are transferred forming ions. Ions are held together due to attractions of opposite charges. (electrostatic attractions) Covalent bonds occur between two nonmetals. The atoms share two or more pair of electrons. Nonpolar bonds have equal sharing of electrons. Polar bonds have an unequal sharing of electrons.

3. Example of Ionic Bonding

4. Ionic Compound Properties 1. Form crystalline solids 2. Have high melting points 3. Conduct electric currents when melted/dissolved

5. Writing an Ionic Formula 1.Write the symbol and charge for the cation. 2.Write the symbol and charge for the anion. 3.Determine subscripts as needed to make total positive charge equal to total negative charge. (Cross and Drop!) Magnesium chloride Potassium carbonate Iron(III) oxide Ammonium sulfate

6. Transition metal ionic compounds  Transition metals can form more than one cation, with different ionic charges.  We indicate the charge on these metals with Roman numerals FeCl 2 the charge on the iron is ? iron(II) chloride FeCl 3 iron(III) chloride Cr 2 S 3 chromium(III) sulfide 2.7

7. Naming Ionic Compounds 1.PbF 2 2.AgNO 3 3.K 2 CO 3 4.Al(OH) 3  Metals (cations) named the same as elements.  Nonmetals (anions) end in –ide  Transition metals need a Roman numeral to indicate charge  Polyatomic ions ( element + oxygen ) end in: –ate (more oxygen), or –ite (less oxygen).

8. Atoms are held together by ‘sharing’ a pair of electrons in a covalent bond Molecule – neutral group of atoms joined by a covalent bond. Covalent (Molecular) Compounds

9. (c) 2006, Mark Rosengarten Examples of Covalent Bonding

10. Covalent compounds Lewis dot diagrams: 1.Count total valence electrons in compound. 2.Connect elements with single bonds (1 pair of electrons). 3.Use remaining electrons to give each atom 8, except H which needs two! NH 3 SF 2 CH 2 I 2

11. VSEPR Model: Used to determine the shape of covalent compounds. Pairs around the “central” atom ShapeModel Example LDD 1 shared pair 3 unshared pair LinearHCl 2 shared pair 2 unshared pair Bent or angularH2OH2O 3 shared pair 1 unshared pair Triangular pyramidNH 3 4 shared pairTetrahedralCH 4

12. Polarity refers to a separation of electric charge within a molecule In most covalent bonds, the electrons are not shared equally If one element is more electronegative than the other, it pulls the shared electrons closer. This gives the element a slight negative charge – and the other element a slight positive charge

13. Electronegativity difference range Most probable type of bond Example 0.0 – 0.4Non-polar covalentH-H (0.0) 0.5 – 2.0polar covalentH – Cl (0.9) >2.0ionicLi – F (3.0) ≥ 2.0IonicNaCl (2.1) Are the following bonds nonpolar covalent, polar covalent, or ionic? a. P-Cl b.H-H c.K-Cl d.H-O Electronegativity table

14. Types of Formulas Molecular FormulaMolecular Formula The formula that states the actual number of each kind of atom found in one molecule of the compound. Empirical FormulaEmpirical Formula The formula of a compound that expresses the smallest whole number ratio of the atoms present. The molecular formula of glucose is C 6 H 12 O 6 The empirical formula of glucose would be CH 2 O.

16. Intermolecular Forces The forces of attraction between two different molecules These forces are weaker than ionic, covalent, or metallic bonds The physical properties of melting point, boiling point, vapor pressure, evaporation, surface tension, and solubility are due to the strength of attractive forces between molecules

17. Types of Intermolecular Forces 1.Van der Waals Forces (the weakest intermolecular attractions) a. Dipole Interactions b. London Dispersion Forces 2.Hydrogen Bonds ( strong intermolecular force )

18. Hydrogen Bonding in Water Hydrogen bonding is an important factor that influences the structure and properties of water One oxygen atom can be hydrogen bonded to as many as 6 other hydrogen atoms